View
218
Download
2
Tags:
Embed Size (px)
Citation preview
Chapter 13
Chemical EquilibriumChemical Equilibrium
Copyright © Houghton Mifflin Company. All rights reserved. 13–2
Figure 13.1: A molecular representation of the reaction 2NO2(g) N2O4(g) over time in a closed
vessel.
Copyright © Houghton Mifflin Company. All rights reserved. 13–3
Figure 13.2: The changes in concentrations with time for the reaction H2O(g) + CO(g) H2(g) + CO2(g)
when equimolar quantities of H2O(g) and CO(g) are mixed.
Copyright © Houghton Mifflin Company. All rights reserved. 13–4
Figure 13.3: (a) H2O and CO are mixed in equal numbers and begin to react (b) to form CO2 and H2. After time has
passed, equilibrium is reached (c) and the numbers of reactant and product molecules then remain
constant over time (d).
Copyright © Houghton Mifflin Company. All rights reserved. 13–5
Figure 13.4: The changes with time in the rates of forward and reverse reactions for H2O(g) + CO(g)H2(g) + CO2(g) when equimolar quantities of H2O(g) and CO(g) are mixed. The rates do not change in the same
way with time because the forward reaction has a much larger rate constant than the reverse reaction.
Copyright © Houghton Mifflin Company. All rights reserved. 13–6
Figure 13.5: A concentration profile for the reaction N2(g) + 3H2(g)2NH3(g) when only
N2(g) and H2(g) are mixed initially.
Copyright © Houghton Mifflin Company. All rights reserved. 13–7
Copyright © Houghton Mifflin Company. All rights reserved. 13–8
Figure 13.6: The position of the equilibrium CaCO3(s) CaO(s) + CO2(g) does not depend on the amounts of CaCO3(s) and
CaO(s) present.
Copyright © Houghton Mifflin Company. All rights reserved. 13–9
Hydrated copper (II) sulfate on the left. Water applied to anhydrous copper (II) sulfate, on
the right, forms the hydrated compound.
Figure 13.7: (a) A physical analogy illustrating the difference between thermodynamic and kinetic stabilities. (b) The reactants H2 and O2 have a strong tendency to form H2O.
Copyright © Houghton Mifflin Company. All rights reserved. 13–11
Apollo II lunar landing module at Tranquility Base, 1969.
Copyright © Houghton Mifflin Company. All rights reserved. 13–12
Procedure for Solving Equilibrium Problems
Copyright © Houghton Mifflin Company. All rights reserved. 13–13
Copyright © Houghton Mifflin Company. All rights reserved. 13–14
Figure 13.8: (a) The initial equilibrium mixture of N2, H2, and NH3. (b) Addition of N2.(c) The new equilibrium position for the system containing more N2 (due to addition of N2),
less H2, and more NH3 than in (a).
Copyright © Houghton Mifflin Company. All rights reserved. 13–15
Figure 13.9: (a) A mixture of NH3(g), N2(g), and H2(g) at equilibrium. (b) The volume is suddenly decreased. (c) The new equilibrium position for
the system containing more NH3 and less N2 and H2. The reaction N2(g) + 3H2(g) 2NH3(g) shifts to the right (toward the side with
fewer molecules) when the container volume is decreased.
Figure 13.10: (a) Brown NO2(g) and colorless N2O4(g) in equilibrium in a syringe. (b) The volume is suddenly decreased, giving a greater concentration of both N2O4 and NO2 (indicated by the darker brown color). (c) A few seconds after the sudden volume decrease, the color is much lighter brown as the equilibrium shifts the brown NO2(g) to colorless N2O4(g) as predicted by Le Châtelier’s principle, since in the equilibrium2NO2(g) ∆ N2O4(g)the product side has the smaller number of molecules.
Copyright © Houghton Mifflin Company. All rights reserved. 13–17
Shifting the N2O4(g) 2NO2(g) equilibrium by changing the temperature. (a) At 100ºC the flask is definitely reddish
brown due to a large amount of NO2 present. (b) At 0ºC the equilibrium is shifted toward colorless
N2O4(g).
Copyright © Houghton Mifflin Company. All rights reserved. 13–18
Copyright © Houghton Mifflin Company. All rights reserved. 13–19